Chemical RelationshipsAcids and Bases: An Introduction

Did you know that some juices and vinegar taste sour because of the chemical properties of the acid in those liquids? And that when acids and bases are mixed together, they always counteract each other, producing water and a salt?

Summary

Since acids and bases were first labeled and described in the 17th century, their definition has been refined over the centuries to reflect an increased understanding of their chemical properties. This module introduces the fundamentals of acid/base chemistry, including neutralization reactions. The relationship between hydrogen ion concentration [H+] and pH is shown alongside everyday examples of acids and bases.

Terms you should know

hydroxide ion = a negatively charged chemical compound that contains one hydrogen atom and one oxygen atom, written as OH-

For thousands of years people have known that vinegar, lemon juice, and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour – because they are all acids. The term acid, in fact, comes from the Latin term acere, which means "sour". While there are many slightly different definitions of acids and bases, in this lesson we will introduce the fundamentals of acid/base chemistry.

In the seventeenth century, the Irish writer and amateur chemist Robert Boyle first labeled substances as either acids or bases (he called bases alkalies), according to the following characteristics:

Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases.

Bases feel slippery, change litmus blue, and become less basic when mixed with acids.

While Boyle and others tried to explain why acids and bases behave
the way they do, the first reasonable definition of acids and bases
would not be proposed until 200 years later.

In the late 1800s, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius
suggested that acids are compounds that contain hydrogen and can
dissolve in water to release hydrogen ions into solution.
For example, hydrochloric acid (HCl) dissolves in water as follows:

HCl

H2O→

H+(aq)

+

Cl-(aq)

Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH):

NaOH

H2O→

Na+(aq)

+

OH-(aq)

The Arrhenius definition of acids and bases explains a number of
things. Arrhenius's theory explains why all acids have similar
properties to each other (and, conversely, why all bases are similar):
because all acids release H+ into solution (and all bases
release OH-). The Arrhenius definition also explains
Boyle's observation that acids and bases counteract each other.
This idea, that a base can make an acid weaker, and vice versa, is
called neutralization.

Neutralization

As you can see from the equations, acids release H+
into solution and bases release OH-. If we were to mix
an acid and base together, the H+ion would combine with the OH-
ion to make the molecule H2O, or plain water:

Though Arrhenius helped explain the fundamentals of acid/base
chemistry, unfortunately his theories have limits. For example,
the Arrhenius definition does not explain why some substances, such as
common baking soda (NaHCO3), can act like a base even though
they do not contain hydroxide ions.

In 1923, the Danish scientist Johannes Brønsted
and the Englishman Thomas Lowry published independent yet similar papers
that refined Arrhenius' theory. In Brønsted's words, "... acids and bases are substances that are capable of splitting off or taking up hydrogen ions, respectively."
The Brønsted-Lowry definition broadened the Arrhenius concept of acids
and bases.

The Brønsted-Lowry definition of acids is very similar to the
Arrhenius definition: Any substance that can donate a hydrogen ion is an
acid. (Under the Brønsted definition, acids are often referred to as proton
donors because an H+ ion, hydrogen minus its electron, is
simply a proton).

The Brønsted definition of bases
is, however, quite different from the Arrhenius definition. The
Brønsted base is defined as any substance that can accept a hydrogen
ion. In essence, a base is the opposite of an acid. NaOH and
KOH, as we saw above, would still be considered bases because they can
accept an H+ from an acid to form water. However, the Brønsted-Lowry definition
also explains
why substances that do not contain OH- can act like bases.
Baking soda (NaHCO3), for example, acts
like a base by accepting a hydrogen ion from an acid as illustrated
below:

Acid

Base

Salt

HCl

+

NaHCO3

→

H2CO3

+

NaCl

In this example, the carbonic acid formed (H2CO3) undergoes rapid decomposition to water and gaseous carbon dioxide, and so the solution bubbles as CO2 gas is released.

pH

Under the Brønsted-Lowry definition, both acids and bases are related
to the concentration of hydrogen ions present. Acids increase the
concentration of hydrogen ions, while bases decrease the concentration
of hydrogen ions (by accepting them). The acidity or basicity of
something, therefore, can be measured by its hydrogen ion concentration.

In 1909, the Danish biochemist Sören Sörensen
invented the pH
scale for measuring acidity. The pH scale is described by the
formula:

pH = -log [H+]

Note: Concentration is commonly abbreviated by using square brackets, thus [H+] = hydrogen ion concentration. When measuring pH, [H+] is in units of moles of H+ per liter of solution.

For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal to 7 (a simpler way to think about pH is that it equals the exponent on the H+ concentration, ignoring the minus sign). The pH scale ranges from 0 to 14. Substances with a pH between 0 and less than 7 are acids (pH and [H+] are inversely related - lower pH means higher [H+]). Substances with a pH greater than 7 and up to 14 are bases (higher pH means lower [H+]). Right in the middle, at pH = 7, are neutral substances, for example, pure water. The relationship between [H+] and pH is shown in the table below alongside some common examples of acids and bases in everyday life.